Group 7

Iodine’s solubility

• It’s not soluble in water.

• However, if iodide ions are present, then it is soluble and reacts to form triiodide ions, which gives the solution it’s brown colour.

• I2 + I- → I3-

Trends down group 7

• Melting and boiling points increase because larger molecules and larger number of v.d.ws.

• Electronegativity decreases down the group because increased atomic radius and shielding.

• Ionisation energy decreases because of shielding and increased atomic radius which weakens the forces of attraction between valence electrons and nucleus.

Chlorine reacting with water

• Reacts with water in a disproportionation reaction, meaning the chlorine is simultaneously oxidised and reduced.

• Cl2 (g) + H2O → ← HCl (aq) + HClO (aq)

• HClO is chloric acid. It’s a mild oxidising agent and effective at killing bacteria without harming humans so it’s used in swimming pools and water treatment for sterilisation.

Chlorine reacting with cold dilute alkali

• Disproportionates again to form chloride, chlorate ion and water.

• Cl2 + 2OH- (aq) → Cl- (aq) + ClO- (aq) + H2O(l)

• The chlorate ion is an important oxidising agent and is used in domestic bleach, NaClO

Appearance and properties of fluorine

• Yellow gas

• Very reactive

• Very toxic

Appearance and properties of chlorine

• Green gas

• Very reactive

• Very toxic

Appearance and properties of bromine

• Brown liquid

• Very easily form orange vapour.

• Orange when dissolved in water.

• Very reactive and toxic.

Appearance and properties of iodine

• Grey crystalline solid/

• Very easily forms purple vapour.

• Solution in water is brown.

2 more reactions of chlorine with water

• 2Cl2 (g) + 2H2O (l) → 4HCl (aq) + O2 (g)

• NaClO (s) + H2O (l) →← Na+ (aq) +OH- (aq) + HClO (aq)

Oxidising power of halogens

• Oxidising agents get reduced easily in order to oxidise another substance, meaning it can gain electrons quickly which they can do because they are very electronegative.

• Oxidising power decreases down a group due to increased shielding and atomic radius.

• Fluorine is the best oxidising agent so it’s ignored because it’s too strong.

• Out of Cl, Br and I, the oxidising power decreases.

Reducing power of the halides

• Reducing agents are oxidised easily, meaning they lose electrons.

• This increases down halides due to decreased electronegativity, making it easier to lose electrons.

• Iodide= fairly good

• Bromide= Fairly poor

• Chloride= Poor

• Fluoride= Very Poor

Identifying halide ions

• Add dilute nitric acid (to ensure carbonates and hydroxides are removed as CO2 or water and do not interfere with the precipitates, producing false positives.

• Add silver nitrates

• Chloride= white

• Bromide= cream

• Iodide= yellow

• Chloride dissolves upon addition of dilute NH3

• Bromide (and chloride) dissolves upon addition of concentrated NH3. (solubility decreases down the group)

Addition of chlorine water to potassium chloride solution

No reaction

Addition of bromine water to potassium chloride solution

No reaction

Addition of iodine solution to potassium chloride solution

No reaction

Addition of chlorine water to potassium bromide solution

• Orange solution as Br2 is made

• Cl2 + 2Br- → 2Cl- +Br2

• Bromide is displaced by chlorine as it’s less reactive than chlorine.

Addition of bromine water to potassium bromide solution

No reaction

Addition of iodine solution (doesn’t dissolve in water) to potassium bromide solution

No reaction

Addition of chlorine and bromine water to potassium iodide

• Brown solution made as I2 is formed.

• This is because iodide is displaced.

Drinking water

Halides as reducing agents: F- and Cl-

• Fluoride and chloride suck and don’t undergo a redox reaction.

• NaCl + H2SO4 → HCl + NaHSO4

• same with fluoride

Bromide as reducing agent

• Lowkey an amazing reducing agent

• First undergoes acid-salt step

• NaBr + H2SO4 → HBr + NaHSO4

• Then redox reaction in which sulfuric acid is reduced to sulfur dioxide.

• 2NaBr + 3H2​SO4​→Br2 ​+ SO2 ​+ 2H2​O + 2NaHSO4​

  • 2Br− → Br2 ​+ 2e−

  • H2​SO4 ​+ 2H+ +2e− → SO2 ​+ 2H2​O

Iodide as reducing agent

• NaI +H2SO4 → HI +NaSO4

• Then redox

• 2NaI + 3H2​SO4​ → I2 ​+ SO2​ + 2H2​O+ 2NaHSO4​

  • 2I−→I2​+2e−

  • H2​SO4​ + 2H+ +2e− → SO2​ + 2H2​O

• 6NaI + 7H2​SO4​ → 3I2 ​+ S + 4H2​O + 6NaHSO4​

  • 6I− → 3I2 ​+ 6e−

  • H2​SO4​ + 6H+ +6e−→S + 4H2​O

• 8NaI + 9H2​SO4 ​→ 4I2 ​+ H2​S + 4H2​O + 8NaHSO4

  • 8I−→4I2​+8e−

  • ​H2​SO4​ + 8H+ +8e− → H2​S + 4H2​O